Pharmaceutical Analytical Chemistry Notes

Pharmaceutical Analytical Chemistry

Analytical chemistry focuses on methods for determining the composition of materials.

• Qualitative analysis: Identifies the components of a material.
• Quantitative analysis: Determines the amounts of each component in a material.

Areas of Chemical Analysis and Questions Addressed

• Identification: What is the identity of the substance in the sample?
• Detection: Does the sample contain a specific substance (X)?
• Quantitation: How much of substance (X) is present in the sample?
• Separation: How can the species of interest be isolated from the sample matrix for better identification and quantitation?

Quantitative Chemical Analysis

• Determines the quantity and purity of a substance.

Classification of Quantitative Methods

A-According to the quantity to be analyzed:

1. Micro methods: For quantities less than 1 mg.
2. Semi-micro methods: For quantities from 1-100 mg.
3. Macro methods: For quantities greater than 100 mg.

B-According to the technique:

1. Volumetric or Titrimetric Methods: Analysis by volume.

2. Gravimetric Methods: Analysis by weight.

3. Instrumental Methods (Physicochemical Methods): Measure a physical property related to the material's concentration.

   • Instrumental methods include:
     • Electrochemical Methods
     • Spectroscopic Methods
     • Separation Techniques

Volumetric or Titrimetric Analysis

• Quantitative chemical analysis by determining the volume of a solution with accurately known concentration needed to react quantitatively with a measured volume of the substance being analyzed.
• Standard solution (Titrant): A solution with an accurately known concentration.
• Titration: The process of gradually adding the standard solution to the sample until the reaction is complete.
• End point or Equivalence point: The point at which the reaction is complete.
• The concentration of the analyzed substance is calculated from the volume of the standard solution used.

Detection of the End Point

• The end point can be determined by:
  1. A physical change produced by the standard solution itself (self-indicator).
  2. Adding an indicator, an organic compound with different colors under different conditions.

Requirements for Quantitative Titrimetric Analysis

1. The reaction must be simple and representable by a chemical equation.
2. The reaction must be instantaneous (or relatively fast); sometimes a catalyst is required.
3. The reaction must occur completely with the titrant in a stoichiometric manner (definite ratio).
4. The end point must be easily detectable (an appropriate indicator is available).

Types of Reactions in Titrimetric Analysis

1. Neutralization Reactions: (Acid-Base reactions)
2. Precipitation Reactions: (Precipitimetry)
3. Complex Formation Reactions: (Complexometry)
4. Electron Transfer Reactions: (Redoximetry)

Standard Solutions

• Solutions of exactly known concentration.

Types of Standard Solutions:

1. Molar Standard Solution (M): Contains the gram molecular weight of the substance in 1 liter of solution.

   • 1M solution: 1 gram molecular weight of substance/L of solution.

   • 2M solution: 2 x gram molecular weight of substance/L of solution.

   • M/10 solution: 0.1 x gram molecular weight of substance/L of solution.

   • Examples:

     • 1 M solution of $H<em>2SO</em>4$ contains 98.07 gm/L of solution.
     • 2 M solution of $H<em>2SO</em>4$ contains 196.14 gm/L of solution.
     • M/10 solution of $H<em>2SO</em>4$ contains 9.8 gm/L of solution.
     • 1 M solution of NaOH contains 40 gm / L of solution.
     • 2 M solution of NaOH contains 80 gm /L of solution.
     • M/10 solution of NaOH contains 4 gm/ L. of solution.
     • 1 M solution of $Na<em>2CO</em>3$ contains 106 gm/L of solution.
     • 2 M solution of $Na<em>2CO</em>3$ contains 212 gm/L of solution.

2. Normal Standard Solution (N): Contains gram equivalent weight/L of solution.

   • 1N solution: 1 gram equivalent weight of substance / L of solution.

   • 2N solution: 2 x gram equivalent weight of substance / L of solution.

   • N/10 solution: 0.1 x gram equivalent weight of substance/L of solution.

   • Equivalent Weight:

     • Acids: Molecular weight / number of replaceable $H^+$
       • Example: Eq. wt of HCl = m.wt/1; Eq. wt of $H<em>2SO</em>4$ = m.wt/2
     • Bases: Molecular weight / number of replaceable $OH^-$
       • Example: Eq.wt of NaOH = m. wt/1; Eq.wt of $Ba(OH)_2$ = m.wt/2
     • Salts:
       • Salts of weak acid and strong base: Eq. wt is equal to the weight which reacts with one $H^+$. Example: $Na<em>2CO</em>3$, Eq. wt = M. wt/2
       • Salts of weak base and strong acid: Eq. wt is equal to the weight which reacts with one $OH^-$. Example: $NH_4Cl$, Eq. wt = M. wt/1
       • In general: Eq. wt of salts is equal to m.wt/(number of metal x its charge).
         • Examples: NaCl, Eq. wt = M. wt/1; $CaCl_2$, Eq. wt = M. wt/2
     • Oxidants and reductants: Eq. wt = m.wt/no of electrons transferred by one molecule.
       • Examples: Eq. wt of $FeSO<em>4$ =m.wt/1; Eq. wt of $CeSO</em>4$ =m.wt/ 2; Eq. wt of $KMnO<em>4$ =m.wt/5; Eq. wt of $K</em>2Cr<em>2O</em>7$=m.wt/6

   • NB: Equal volumes of equi-normal solutions contain equal numbers of molecules, meaning equal normalities react in a 1:1 ratio.

3. Empirical Standard Solution:

   • Prepared such that 1 ml reacts with a definite quantity of another substance.
   • Specific to the determination of certain substances and not useful for others (no relation between different empirical solutions).
   • Not common in analytical chemistry but used mainly in industry.

Neutralization Reactions: Acid-Base Titrations in Aqueous Solutions

Solutions

• Homogeneous mixture of two or more substances.

  • Solute: The component present in a small quantity (solid, gas, or liquid).
  • Solvent: The component present in a large quantity.

• Solutions may be:

  1. Saturated Solutions
  2. Unsaturated Solutions
  3. Supersaturated Solutions

Electrolytes and Non-electrolytes

• Electrolytes: Substances that dissociate into ions when dissolved in water, creating a solution that conducts electricity.
  • Strong Electrolytes: Dissociate or ionize to a high degree in water.
    • Examples:
      • Acids: HCl, $HNO<em>3$, $H</em>2SO_4$, HBr, HI.
      • Bases: NaOH, KOH, $Ca(OH)<em>2$, $Ba(OH)</em>2$.
      • Salts: NaCl, $CH<em>3COONa$, $NH</em>4Cl$.
  • Weak Electrolytes: Ionize to a slight degree.
    • Examples:
      • Acids: $CH<em>3COOH$, $HBO</em>3$, HF, HCN, $H_2S$.
      • Bases: $NH<em>4OH$, $N</em>2H_4$.
      • Salts: $HgCl<em>2$, $CdCl</em>2$,$HgBr<em>2$, $CH</em>3COONH_4$.
• Non-Electrolytes: Substances that do not ionize when dissolved in water and yield non-conducting solutions (e.g., sugar, glycerin, ethyl acetate).

Electrolytic Dissociation Theory

• Pure water is a poor conductor of electricity.
• When an electrolyte dissolves in water, it dissociates into negatively charged ions (anions) and positively charged ions (cations).
• The solution remains electrically neutral, with the total number of negative charges equaling the total number of positive charges.
• Acidic properties are due to hydrogen ions ($H^+$), and basic properties are due to hydroxide ions ($OH^-$).
• Solutions conduct electric current due to the presence of ions.
• Degree of Dissociation ($\alpha$):
  • The ratio of ionized fraction to the total amount of dissolved solute.
  • For each concentration, there is an equilibrium between undissociated molecules and dissociated ions.
  • $Molecule \rightleftharpoons cation (+ve) + Anion (-ve)$
  • $CH<em>3COOH \rightleftharpoons H^+ + CH</em>3COO^-$
  • $NH<em>4Cl \rightleftharpoons NH</em>4^+ + Cl^-$
  • The degree of dissociation characterizes the chemical activity of substances.

Molecular and Ionic Equations

• Molecular Equations: Represent reactants and products as molecules.

  • $NaOH + HCl \rightarrow NaCl + H_2O$
  • Indicates that 1 mole of NaOH neutralizes exactly 1 mole of HCl to form 1 mole of NaCl and 1 mole of $H_2O$.

• Ionic Equations: Strong electrolytes are represented as ions, while weak electrolytes are represented as molecules.

  • For the reaction of NaOH & HCl (both strong electrolytes):
    • $Na^+ + OH^- + H^+ + Cl^- \rightarrow Na^+ + Cl^- + H_2O$
    • Simplified to: $OH^- + H^+ \rightarrow H_2O$
  • In the reaction of NaOH (strong electrolyte) with $CH_3COOH$ (weak electrolyte):
    • $NaOH + CH<em>3COOH \rightarrow Na^+ + CH</em>3COO^- + H_2O$
    • Simplified to:
      • $OH^- + CH<em>3COOH \rightarrow CH</em>3COO^- + H_2O$

Chemical Equilibrium

• In reversible reactions, products are formed from reactants, and reactants are produced from products simultaneously indicated by the symbol ($\rightleftharpoons$).
• $A + B \rightleftharpoons C + D$
• The reaction mixture achieves a constant composition, and the system is in a state of equilibrium.
• Equilibrium is the state where the rate of the forward reaction equals the rate of the backward reaction.

Law of Mass Action

• The rate of a chemical reaction is directly proportional to the product of the molar concentrations of the reacting substances.
• For the reaction: $A + B \rightleftharpoons C + D$
  • $V<em>1 \propto [A][B]$ or $V</em>2 \propto [C][D]$
  • At equilibrium, $V<em>1 = V</em>2$
  • $V<em>1 = K</em>1[A][B]$ or $V<em>2 = K</em>2[C][D]$
  • $K<em>1[A][B] = K</em>2[C][D]$
  • $K<em>{eq} = \frac{K</em>1}{K_2} = \frac{[C][D]}{[A][B]}$
• In the general case:
  • $aA + bB \rightleftharpoons cC + dD$
  • $K_{eq} = \frac{[C]^c[D]^d}{[A]^a[B]^b}$

Displacement of Equilibrium: Le-Chatelier Principle

• If a stress is applied to a system in equilibrium, the equilibrium will shift in a direction that relieves the stress.

Applications of Le-Chatelier Principle

• In Precipitation:

  • $A + B \rightleftharpoons AB_{precipitate}$
  • Adding excess precipitating agent B will cause the system to absorb the excess B by combining with A to form more AB, shifting the equilibrium to the right.

• In Solubility:

  • Endothermic Solution: Heating increases the solubility of the solute (equilibrium shifts to the right).
    • Solute + Solvent + \Delta \rightleftharpoonsSolution
  • Exothermic Solution: Heating decreases the solubility of the solute (equilibrium shifts to the left).
    • $Solute + Solvent \rightleftharpoons Solution + \Delta$

Dissociation of Water and Ionic Product of Water ($K_w$)

• $H_2O \rightleftharpoons H^+ + OH^-$
• Water molecules ionize to a very slight degree.
• According to the law of mass action:
  • $K = \frac{[H^+][OH^-]}{[H_2O]}$
  • $K[H_2O] = [H^+][OH^-]$
  • Since the fraction of water ionized is very small ($10^{-7}$ gm-ion/L), the concentration of water is regarded as approximately constant.
  • $K_w = [H^+][OH^-]$
• $K_w$ is the ionic product of water.
• Under normal conditions and at approximately 25°C:
  • $K_w = [H^+][OH^-] = 10^{-14}$
• The dissociation of water produces equal numbers of $H^+$ and $OH^-$ ions.
• $[H^+] = \sqrt{10^{-14}} = 10^{-7}$

The Hydrogen Ion Exponent (pH)

• pH is a measure of the acidity or alkalinity of a solution.
• pH is the negative logarithm of the hydrogen ion molar concentration.
• $pH = -log[H^+]$
• $pH = n$ if $[H^+] = 10^{-n}$
• Suppose that, the $[H^+] = 10^{-8}$
• $pH = 8$
• $K_w = 10^{-14}$
• $pK_w=14$
• pH range : Acidity increased (0-6), neutral(7), Alkalinity increased (8-14)
• In general:
  • $pH = -log[H^+]$
  • $pOH = -log[OH^-]$
  • $pK_w = pH + pOH$
  • $pH = pK_w - pOH$
  • $pH = 14 - pOH$

pH of Acids and Bases

• pH of strong acids and strong bases:

  • Strong acids and strong bases are completely ionized so the concentration of the acid or base represents the concentration of $H^+$ or $OH^-$.
  • For acids: $pH = -log[H^+]$
  • For bases: $pOH = -log[OH^-]$
  • $pH = 14 - pOH$
  • Examples:
    • pH of 0.1M HCl (strong acid):
      • $pH = -log[H^+] = -log[10^{-1}] = 1$
    • pH of 0.1M NaOH (strong base):
      • $pOH = -log[OH^-] = -log[10^{-1}] = 1$
      • $pH = 14 - pOH = 14 - 1 = 13$

• pH of weak acids: A small quantity of weak acids (e.g., $CH_3COOH$) is dissociated with the formation of $H^+$.

  • $CH<em>3COOH \rightleftharpoons CH</em>3COO^- + H^+$
  • $K<em>a = \frac{[CH</em>3COO^-][H^+]}{[CH_3COOH]}$
  • Where: $[H^+]=[CH<em>3COO^-]$ & $[CH</em>3COOH] = C_a$ (concentration of acid)
  • $K<em>a = \frac{[H^+]^2}{C</em>a}$

• pH of weak bases: As weak acids, weak bases eg NH,OH
  dissociated with the formation of $OH^-$

  • $[OH^-]^2 = K<em>bC</em>b$
  • $[OH^-] = \sqrt{K<em>bC</em>b}$
  • $pOH = pK<em>b+ \frac{1}{2} pC</em>b$
  • $pH = pK_w-pOH$